1. Unit 3 Atomic Structure, Periodic Table and bonding Chapters 7, 8 and 9

2. Homework Due with next test Pg 329 3-7, 21-51,77-103 odd Pg 393 21-31, 41-51, 63-69, 81-91 odd Pg 431 5, 23-29 odd

3. Dalton’s Model (1803) people started to accept the idea of atoms because of his experiments He worked with gases and found that they acted as though they were made of solid microscopic particles all elements are made of atoms (indivisible and indestructible) He took the name atom off of Greek philosopher Democritus’ idea that matter had a smallest indestructible piece to it.

4. John Dalton

5. Thomson’s Cathode Ray

6. Thomson’s Model (1897) found negative particles could come from neutral elements. Therefore he discovered the electrons around the atom. atom is made of smaller things (+ & -), which we now call subatomic particles successfully separated negative particles (electrons) but could not separate the positive particle (protons) “plum pudding model” negative particles floating in a positively charged gel like material

7. Plum Pudding Model- Thomson Positive Gel Negative Particles

8. Sir J. J. Thomson The sir means he was knighted

9. Rutherford’s Model (1911) fired protons at a sheet of gold foil most went through unaffected, some bounced away there is a small dense area of positive particles at the center of the atom- the nucleus He discovered the nucleus of an atom electrons are scattered near the outside of the atom with mostly empty space between the nucleus and the electrons

10. Gold Foil Experiment Radioactive source Gold foil

11. Rutherford’s Model nucleus (small dense positive area) electrons Empty Space

12. Ernest Rutherford

13. Neils Bohr Bohr saw, when he dispersed (put it through a prism to separate it) light from a hydrogen light bulb, distinct bands of colors instead of a smooth transition. What you would expect What he actually saw

14. Hydrogen light bulb A hydrogen light bulb works like a neon light. Electricity goes through the gas. All the atoms jump to an excited state. When they fall back to ground state, they give off energy as light. Since we only see bands of light (when dispersed), we know energy is coming out in discrete amounts, not a steady flow.

15. Energy Levels From this, Bohr determined electrons were at certain energy levels from the nucleus. Excited e - ’s jump to higher energy levels, then fall back to ground state. Since the distance it “falls” back is always the same, energy always comes out of an atom in discrete amounts.

16. Bohr Model (1913) electrons move in definite orbits around the nucleus these orbits or energy levels are located at certain distances from the nucleus He calculated the distance these energy levels should be and made them perfect circles about the atom.

17. Bohr’s Model nucleus Electrons

18. Neils Bohr

19. Wave Model (present day) Bohr based his model off of classical physics. Quantum mechanics was now being developed. The math does not work with classical physics, however the predictions made by quantum mechanics always work. de Broglie stated that electrons (particles) have wave properties, and he viewed these as standing waves, like those produced when a guitar string is plucked (classical physics.) Schrodinger assumed that the electron in Hydrogen behaves as a standing wave.

21. What is an orbital? It is not a Bohr orbit (not moving in a circular path.) How is the electron moving? We don’t know! There is a fundamental limitation to just how precisely we can know both the position and momentum of a particle at a given time

22. This is kind of how we assume an electron travels e-e-

23. The Heisenberg Uncertainty Principle The Heisenberg uncertainty principle states that it is impossible to know simultaneously the exact position and velocity of a particle. Δx = uncertainty in a particle’s position Δ(mν) = uncertainty in a particle’s momentum h = Planck’s constant

24. The Quantum-Mechanical Model of the Atom The Schrodinger equation is an extremely complex equation used to describe the 3-D quantum mechanical model of the hydrogen atom. The hydrogen electron is visualized as a standing wave (a stationary wave as one found on a musical instrument) around the nucleus. The circumference of a particular orbit would have to correspond to a whole number of wavelengths.

25. This is consistent with the fact that only certain electron energies are shown to exist.

26. The Physical Meaning of a Wave Function The square of the function indicates the probability of finding an electron near a particular point in space. The probability distribution, or electron density diagram, can be plotted to give the most probable region where an electron may be located.

27. Solving Schrodinger’s Equations Schrodinger’s equations give us different values for the orbitals in a hydrogen atom. These values are called quantum numbers n the principle quantum number. This relates to the size and energy of the orbital. l the angular momentum. This relates to the shape of the orbital. m l the magnetic quantum number. This relates to the orientation of the orbital in space to the other orbitals in the atom

28. Quantum Numbers for the First Four Levels of Orbitals in the Hydrogen Atom.

29. Orbital diagrams From these quantum numbers, orbital diagrams were designed. 1s_ 2s_p_ _ _ 3s_p_ _ _ d _ _ _ _ _ 4s_p_ _ _ d _ _ _ _ _ f _ _ _ _ _ _ _

30. Orbital shapes Shapes are denoted by letters s, p, d, f, g … s orbitals have a spherical shape Separating s orbitals (the space between 1s and 2s) there are areas of 0 probability of finding an electron called nodes. p orbitals are pear shaped with a node in between them http://orbitals.com/orb/

31. Reviewing orbitals All are based on hydrogen atom Electrons are viewed as a standing wave. We don’t know how electrons move, instead we represent their locations with probability distribution maps. Size of the orbital is defined at the area that contains 90% of the probable location of the atom The ground state for hydrogen is 1s orbital

32. Louis de Broglie Erwin Schrodinger

33. Werner Heisenberg

34. Electron Spin A 4 th quantum number was discovered to be necessary to account for emission spectra of atoms. Electrons appeared to have a magnetic moment (North pole and South pole) that pointed in one of two orientations. Since looping or spinning electric current creates a magnetic field, this was called the spin of the electron m s is electron spin, it can be +½ or –½

35. Pauli Exclusion Principle In a given atom, no two electrons may have the same set of four quantum numbers. This leads us to there may only be two electron in one orbital, and they must have an opposite spin value. The electrons of an atom in its ground state occupy the orbitals of lowest energy.

36. Electrostatic Effects In one electron systems, the only electrostatic force is the nucleus-electron attractive force. Many-electron systems (polyelectronic atoms). Nucleus-electron attractions. Electron-electron repulsions.

37. Polyelectronic Atoms For polyelectronic atoms, one major consequence of having attractions and repulsions is the splitting of energy levels into sublevels of differing energies. The energy of an orbital in a many-electron atom depends on its n value (size) and secondly on its l value (shape).

38. This means that energy levels may split into s, p, d, and f sublevels, depending on the n value.

39. Copyright © Cengage Learning. All rights reserved 39 Orbital Energies

40. Electrostatic Effects In one electron systems, the only electrostatic force is the nucleus-electron attractive force. Many-electron systems (polyelectronic atoms). Nucleus-electron attractions. Electron-electron repulsions.

41. Polyelectronic Atoms For polyelectronic atoms, one major consequence of having attractions and repulsions is the splitting of energy levels into sublevels of differing energies. The energy of an orbital in a many-electron atom depends on its n value (size) and secondly on its l value (shape).

42. This means that energy levels may split into s, p, d, and f sublevels, depending on the n value.

43. Copyright © Cengage Learning. All rights reserved 43 Orbital Energies

44. Electrostatic Effects In one electron systems, the only electrostatic force is the nucleus-electron attractive force. Many-electron systems (polyelectronic atoms). Nucleus-electron attractions. Electron-electron repulsions.

45. Polyelectronic Atoms For polyelectronic atoms, one major consequence of having attractions and repulsions is the splitting of energy levels into sublevels of differing energies. The energy of an orbital in a many-electron atom depends on its n value (size) and secondly on its l value (shape).

46. This means that energy levels may split into s, p, d, and f sublevels, depending on the n value.

47. Copyright © Cengage Learning. All rights reserved 47 Orbital Energies

48. Conclusions A greater nuclear charge (Z) lowers orbital energies. Why? Electron-electron repulsions raise orbital energy. Why? Electrons in outer orbitals (higher n) are higher in energy. Why?

49. Periodic Table

50. History of the periodic table. Johann Dobereiner (1780-1849). first known chemist to recognize periodic patterns. suggested a “triad” system for classification, but found it limiting. John Newlands. * In 1864, he suggested that elements should be arranged in “octaves,” based on the idea that certain properties seemed to repeat every eighth element; still not very successful.

51. Atomic Mass Avogadro had suggested the idea of a mole. Which was a specific number of atoms, he had no idea of actual the amount, 6.022x10 23. He said equal volumes of gases at the same pressure and temperature would have the same number of particles, thus 1 mole of any gas at STP = 22.4 L By reacting moles of elements with each other you can determine what one mole of every element is. If you take the mass of one mole of an element you have the atomic mass.

52. Mendeleev and Meyer Julius Meyer and Dmitri Mendeleev both independently conceived the present periodic table. Mendeleev is given credit because of his ability to predict accurately the properties of undiscovered elements.

53. Dmitri Mendeleev Element #101 was named in his honor first to come up with a periodic table ~1870 there were 63 elements known to man, he organized the elements in order of their atomic mass, he saw a pattern form with the # of elements that can be bonded to that element.

54. More Mendeleev He arranged his table so that columns of elements with repeating patterns formed, he left spaces in his table where it appeared an element wasn’t discovered yet He then predicted the properties of these elements by looking at the other elements in the family and period of that element

55. Medeleev’s Table

56. It may not sound like much, but… predictedundiscoveredHe predicted what undiscovered elements would be like giving specifics!!! Right!And he was Right! This is the basis for acceptance of a scientific theory. If the model is correct then the scientist can assume what would happen if…

57. 1871 Discovery PropertyEkaaluminum (Mendeleev’s prediction) Gallium atomic mass6869.72 density (g/cm³)6.05.904 melting point (°C) Low29.78 oxide's formulaEa 2 O 3 (density - 5.5 g cm -3 ) (soluble in both alkalis and acids) Ga 2 O 3 (density - 5.88 g cm -3 ) (soluble in both alkalis and acids) chloride's formula Ea 2 Cl 6 (volatile)Ga 2 Cl 6 (volatile)

58. 1882 Discovery PropertyEkasiliconGermanium atomic mass7272.61 density (g/cm³)5.55.35 melting point (°C)high947 colorgrey oxide typerefractory dioxide oxide density (g/cm³)4.7 oxide activityfeebly basic chloride boiling pointunder 100°C86°C (GeCl 4 ) chloride density (g/cm³)1.9

59. Dmitri Mendeleev

60. Problem with Mendeleev’s table Mendeleev arranged his table by atomic mass a few elements appeared to be slightly out of place, Mendeleev put them in the right place and guessed that their atomic masses were incorrectly measured Actually, he was arranging them by the wrong number.

61. Atomic Number and Henry Moseley Moseley discovered that the charge of a nucleus is always divisible by the same number (which we now call of the charge of a proton). This gave him an integral number for each atom he called the atomic number. He corrected the periodic law when he stated that periodicity is a result of increasing atomic number, not increasing atomic mass.

62. Henry Moseley